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Comparing formal charges to oxidation states

How formal charges and oxidation states are both ways of counting electrons.

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  • marcimus pink style avatar for user Moke Thon Hein
    Why are the definitions "hypothetical"? Are formal charges solely hypothetical or are they practical?
    (11 votes)
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    • piceratops tree style avatar for user Christopher
      The most useful definition of hypothetical for you would be : "supposed but not necessarily real or true."

      In this instance, the hypothetical charge is the charge calculated by assuming either a) bonding electrons are shared equally or b) bonding electrons are assigned to the more electronegative atom

      Therefore, these are hypothetical values because they are only true after we have made some assumption. They are not intrinsic to the atoms themselves.
      (24 votes)
  • mr pants teal style avatar for user sama.elberashy
    what is the use of formal charge or oxidation number??
    (13 votes)
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  • piceratops tree style avatar for user Christopher
    What would be the formal charge/oxidation number for the oxygen in these exercises. I paused the video and tried to work it out. Here is my reasoning:

    Formal charge - Oxygen has six valence electrons and two bonds. So the formal charge would be 6 - 2 = 4

    Oxidation state - Oxygen has six valence electrons and two bonds. It is the more electronegatative element for both bonds. Therefore, it's oxidation state would be 6 - 2 - 2 = 2

    Is this correct?
    (4 votes)
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  • female robot amelia style avatar for user Tejaswini
    does the free end of the structure always represent methyl group?
    (5 votes)
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  • leafers sapling style avatar for user Luxmy
    What is the use of formal charge and oxidation charge? I understand how to do it, but I don't know why we need to do it.
    (7 votes)
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  • blobby green style avatar for user Murad
    why do we need concepts like oxidation states, and formal charge.?
    (5 votes)
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  • old spice man blue style avatar for user Lucian Rex
    Since alcohols have a OH bonded to it are they basic?
    (3 votes)
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  • starky ultimate style avatar for user alina
    At he says "methanol." I know this is a pretty elementary question, but why does methanol end in "ol"? As for a question actually pertaining to the video, does anyone have classic examples of molecules being used to teach oxidation states. (like practice problems?)
    (2 votes)
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  • blobby green style avatar for user Priya
    can we see carbon structures by any equipment or are they imaginary structures?
    (2 votes)
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  • winston baby style avatar for user Gwen Teng
    what do mean by "the more electronegative atom" ?
    (1 vote)
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Video transcript

- Both formal charge and oxidation states are ways of counting electrons, and they're both very useful concepts. Let's start with formal charge. So one definition for formal charge is the hypothetical charge that would result if all bonding electrons are shared equally. So let's go down to the dot structure on the left here, which is a dot structure for methanol, and let's assign a formal charge to carbon. We need to think about the bonding electrons or the electrons in those bonds around carbon, and we know that each bond consists of two electrons. So the bond between oxygen and carbon consists of two electrons. Let me go ahead and draw in those two electrons. Same for the bond between carbon and hydrogen, right? Each bond consists of two electrons, so I can go around and put in all of my bonding electrons. So if we want to assign a formal charge to carbon, we need to think about the number of valence electrons in the free atom or the number of valence electrons that carbon is supposed to have. We already know that carbon is supposed to have four valence electrons, so I could put a four here, and from that four we're going to subtract the number of valence electrons in the bonded atom or the number of valence electrons that carbon has around it in our drawing. And since we're doing formal charge, we need to think about all those bonding electrons being shared equally. So we think about a covalent bond. So if we have two electrons and one bond, and those two electrons are shared equally, we could split them up. We could give one electron to oxygen and one electron to carbon in that bond. We go over here to this carbon hydrogen bond, and we could do the same thing. We have two electrons. We could split up those two electrons. We could give one to carbon and one to hydrogen, and we go all the way around, and we do the same thing over here. Split up those electrons and the same thing here. So how many valence electrons do we see around carbon now? So let me go ahead and highlight them. There's one, two, three, and four. So that's the number of valence electrons around carbon in our drawing. So four minus four is equal to zero. So zero is the formal charge of carbon. So let me go ahead and highlight that here. So in this molecule the formal charge for carbon is zero. Now let's move on to oxidation states, right? So you could also call these oxidation numbers. So one definition for an oxidation state is the hypothetical charge that would result if all those bonding electrons are assigned to the more electronegative atom in the bond. So let's go to the dot structure on the right of methanol and let's assign an oxidation state to that carbon. We need to think about our bonding electrons again, so let's go ahead and put those in, all right? So we know that each bond consists of two electrons. So I'm putting in the two electrons in each bond, and let's think about the oxidation state of that carbon. Well first, we need to know the number of valence electrons in the free atom. So just like before, we know that carbon is supposed to have four valence electrons. So this would be a four, and from that we subtract the number of valence electrons in the bonded atom or the number of valence electrons that carbon actually has in the drawing. This time we need to think about an ionic bond, so we're going to pretend like a covalent bond is an ionic bond, because we're going to assign all of the bonding electrons to the more electronegative atom. So there's no more sharing here. Winner takes all. The more electronegative atom is going to get all of the electrons. So let's think about the electronegativities of carbon versus oxygen, all right? We know that oxygen is more electronegative than carbon. So oxygen takes both of those electrons in that bond. So oxygen gets both of those electrons. Next let's think about the electronegativities of carbon and hydrogen. We know that carbon is a little bit more electronegative than hydrogen. So for these two electrons, carbon's going to take both of them since carbon is more electronegative than hydrogen, and the same thing for our other carbon hydrogen bonds. Carbon is more electronegative than hydrogen, so carbon takes those. Carbon is more electronegative than hydrogen, so carbon takes those. And so, how many electrons do we have around carbon now? Let's count them up. That's one, two, three, four, five, and six. So now we have six electrons around carbon. So four minus six gives us negative two. So here in this example carbon has an oxidation state of negative two. So there's no more sharing when you're doing oxidation states, right? Think about the more electronegative atom and assign both electrons to the more electronegative atom. Both formal charge and oxidation states are just really extreme methods of electron bookkeeping, right? They're not perfect. They're certainly not perfect, right? We're assuming that the electrons are either shared equally, perfectly, or that one atom takes both electrons, and neither of those concepts is perfect in the real world, but it works when we're drawing our dot structures and we're thinking about chemical reactions.